A 500-mL aqueous solution of NaCl was electrolyzed for 15.0 minutes, producing Cl2 gas at the anode. If the pH of the final solution was 12.15, calculate the average current used. (hint: what products are produced at the cathode?)
A. 3.8 A
B. 2.2 A
C. 0.47 A
D. 1.4 A
E. 0.76 A
See your other post and response by Bob Pursley
To calculate the average current used, we need to know the amount of charge that passed through the electrolysis cell.
First, we need to determine the number of moles of Cl2 gas produced at the anode. Since electrolysis of NaCl produces Cl2 gas at the anode and NaOH and H2 gas at the cathode, we can use the stoichiometry of the reaction to calculate the number of moles of Cl2 produced:
2 mol e- + 2 mol Cl- -> 1 mol Cl2
Since each mole of Cl2 requires 2 moles of electrons, the number of moles of Cl2 is equal to the number of moles of electrons transferred.
The charge (Q) passing through the electrolysis cell can be calculated using Faraday's law:
Q = nF
Where:
Q is the charge in coulombs
n is the number of moles of electrons transferred
F is Faraday's constant (96500 C/mol e-)
To calculate the number of moles of Cl2, we need to know the concentration of NaCl in the solution and convert it to moles:
Concentration (M) = moles of solute (NaCl) / volume of solution (L)
Given that we have a 500 mL solution of NaCl, we need to convert it to liters:
500 mL = 0.5 L
Next, we need to calculate the number of moles of NaCl using its molarity. Since we don't have the molarity provided in the question, we cannot proceed with the calculation.
Therefore, we cannot determine the average current used in this electrolysis process based on the given information.